4.4 Energy, Entropy and Equilibrium

Entropy- the number of ways molecules can be arranged, and the number of ways quanta of energy can be arranged

Solids have low entropies, as they have a regular lattice structure where there is little disorder. Gases have the highest entropies, as the molecules are randomly arranged, and they occupy a greater volume with more randomness in their distribution

Molecules with heavier atoms and more atoms have higher entropies

When calculating total entropy change of a reaction you need to consider the entropy changes of the system and the surroundings

For a spontaneous process ΔStotalis positive

For an equilibrium ΔStotalis 0

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4.5 Energy Changes in Solutions

Lattice enthalpy, ΔHLE- the enthalpy change when 1 mole of solid is formed from separate ions. Always large, negative quantities. Become more negative when ionic charges increase and ionic radii decrease, as ions will attract each other more strongly if closer together

Enthalpy of Hydration, ΔHhyd- The enthalpy change for the formation of a solution of ions from 1 mole of gaseous ions.

Enthalpies of hydration are always -ve. The most exothermic values occur for the ions with the greatest charge and the smallest radii

When dealing with other solvents apart from water,Enthalpy of Solvation, ΔHsolvis used

The difference between the enthalpies of hydration of the ions and the lattice enthalpy gives the Enthalpy Change of Solution, ΔHsolution

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5.5 Hydrogen Bonding and Water

Water, ammonia, methane and hydrogen fluoride have higher than expected boiling points and enthalpy changes of vaporisation. This is because they have the ability to form hydrogen bonds.

Hydrogen bonding holds molecules together strongly, so more energy is needed to overcome them before boiling or vaporisation can occur

The greater the electronegativity of the atoms, the stronger the intermolecular bonds, and the higher the boiling point and the enthalpy of vaporisation.

Specific heat capacity- the energy needed to raise 1g of a substance through 1K. Water has an unusually high shc due to hydrogen bonding. A lot of energy is used to overcome hydrogen bonding in clusters of water molecules, so this energy is not available to increase the KE of the molecules.

Unlike most liquids, the density of water decreases when it freezes. In ice, the arrangement of water molecules maximises the hydrogen bonding between them, leading to an open structure with large spaces in it. Therefore the density of ice is lower than the density of water.

When ice melts, the structure collapses and water molecules fall into some of the open spaces, giving water a greater density

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8.2 Strong and Weak Acids and pH

Strong acids have a strong tendency to donate H+, donation is essentially complete. Weak acids have a weaker tendency to donate H+ and the reaction with water is incomplete.

pH= -lg[H+]

For a strong acid, [H+] is equal to the amount of acid [HA] put into the solution, as the reaction goes to completion. Therefore, pH= -lg[H+]

For a weak acid HA H+ + A-. The equilibrium is the acidity constant, Ka. To work out the pH, you need to find 2 assumptions.

Assumption 1- [H+] = [A-]. Equal amount of H+ and A- are formed from HA, but water can produce H+. However, water produces far less H+ than most weak acids, so the ionisation of water can be neglected.

Assumption 2- The amount of HA at equilibrium is equal to the amount of HA put into the solution. We can neglect the fraction of HA which has lost H+

Concentration is a measure of the amount of substance in a given volume of solution. Strength is a measure of the extent to which an acid can donate H+.

Water does ionise slightly. Kw = [H+] [OH-]. Kw is the ionic product of water.